As used herein, the term “supersorbent” shall mean a sorbent as taught in U.S. Pat. No. 5,779,464 entitled “Calcium Carbonate Sorbent and Methods of Making and Using Same”, the teachings of which are hereby incorporated by reference.
As used herein, the term “microporous” shall mean a pore size distribution of less than 5 nanometers. As used herein, the term “mesoporous” shall mean a pore size distribution of from about 5 nanometers to about 20 nanometers.
Atmospheric CO2 concentration has been increasing steadily since the industrial revolution. It has been widely accepted that the while the CO2 concentration was about 280 ppm before the industrial revolution, it has increased from 315 ppmv in 1959 to 370 ppmv in 2001 [Keeling, C. D. and T. P. Whorf. 2002. Atmospheric CO2 records from sites in the SIO air sampling network. In Trends: A Compendium of Data on Global Change. Carbon Dioxide Information Analysis Center, Oak Ridge National Laboratory, U.S. Department of Energy, Oak Ridge, Tenn., U.S.A. This data is also available from http://cdiac.esd.ornl.gov/ftp/maunaloa-co2/maunaloa.co2]. Rising CO2 concentrations has been reported to account for half of the greenhouse effect that causes global warming [IPCC Working Group I. IPCC Climate Change 1995—The Science of Climate Change: The Second Assessment Report of the Intergovernmental Panel on Climate Change; Houghton, J. T., Meira Filho, L. G., Callander, B. A., Harris, N., Kattenberg, A., Maskell K, Eds.; Cambridge University Press: Cambridge, U.K., 1996]. Although the anthropogenic CO2 emissions are small compared to the amount of CO2 exchanged in the natural cycles, the discrepancy between the long life of CO2 in the atmosphere (50-200 years) and the slow rate of natural CO2 sequestration processes leads to CO2 build up in the atmosphere. The IPCC (Intergovernmental Panel on Climate Change) opines that “the balance of evidence suggests a discernible human influence on the global climate.” Therefore, it is necessary to develop cost effective CO2 management schemes to curb its emission.
Many of the envisaged CO2 management schemes consist of three parts—separation, transportation and sequestration of CO2 [FETC Carbon Sequestration R&D Program Plan: FY 1999-2000. National Energy Technology Laboratory, Department of Energy, Washington, D.C., 1999]. The cost of separation and compression of CO2 to 110 bar (for transportation of CO2 in liquid state) is estimated at $30-50 per ton CO2, and transportation and sequestration would cost about $1-3 per ton per 100 km and $1-3 per ton of CO2, respectively [Wallace, D. Capture and Storage of CO2. What Needs To Be Done. Presented at the 6th Conference of the Parties, COP 6, to the United Nations Framework Convention on Climate Change; The Hague, The Netherlands, Nov. 13-24, 2000; www.iea.org/envissu/index.htm]. The capture of CO2 imposes severe energy penalties thereby reducing the net electricity output by as much as 13-37% [Herzog, H.; Drake, E.; Adams, E. CO2 Capture, Reuse, and Storage Technologies for Mitigating Global Climate Change. A White Paper; Final Report No. DE-AF22-96PC1257, January 1997]. The dominating costs associated with the current CO2 separation technologies necessitate development of economical alternatives.
Historically, CO2 separation was motivated by enhanced oil recovery [Kaplan, L. J. Cost-Saving Processes Recovers CO2 from Power-Plant Flue gas. Chem. Eng. 1982, 89 (24), 30-31; Pauley, C. P.; Smiskey, P. L.; Haigh, S. N—ReN Recovers CO2 from Flue Gas Economically. Oil Gas J. 1984, 82(20), 87-92]. Currently, industrial processes such as limestone calcination, synthesis of ammonia and hydrogen production require CO2 separation. Absorption processes employ physical and chemical solvents such as Selexol and Rectisol, MEA and KS-2 [Reimer, P.; Audus, H.; Smith, A. Carbon Dioxide Capture from Power Stations. IEA Greenhouse R&D Programme, www.ieagreen.org.uk, 2001. ISBN 1 898373 15 9; Blauwhoff, P. M. M.; Versteeg, G. F.; van Swaaij, W. P. M. A study on the reaction between CO2 and alkanoamines in aqueous solution. Chem. Eng. Sci.1984, 39(2), 207-225. Mimura, T.; Simayoshi, H.; Suda, T.; Iijima, M.; Mitsuake, S. Development of Energy Saving Technology for Flue Gas Carbon Dioxide Recovery by Chemical Absorption Method and Steam System in Power Plant. Energy Convers. Mgmt. 1997, 38, Suppl. P.S57-S62]. Adsorption systems capture CO2 on a bed of adsorbent materials such as molecular sieves and activated carbon [Kikkinides, E. S.; Yang, R. T.; Cho, S. H. Concentration and Recovery of CO2 from flue gas by pressure swing adsorption. Ind. Eng. Chem. Res. 1993, 32, 2714-2720]. CO2 can also be separated from the other gases by condensing it out at cryogenic temperatures. Polymers, metals such as palladium, and molecular sieves are being evaluated for membrane based separation processes [Reimer, P.; Audus, H.; Smith, A. Carbon Dioxide Capture from Power Stations. IEA Greenhouse R&D Programme, www.ieagreen.org.uk, 2001. ISBN 1 898373 15 9].
Reaction based processes, as promulgated in this work, can be applied to separate CO2 from gas mixtures. This process is based on a heterogeneous gas-solid non-catalytic carbonation reaction where gaseous CO2 reacts with solid metal oxide (represented by MO) to yield the metal carbonate (MCO3). The reaction can be represented by:MO+CO2→MCO3  (1)Once the metal oxide has reached its ultimate conversion, it can be thermally regenerated to the metal oxide and CO2 by the calcination of the metal carbonate product. The calcination reaction can be represented by:MCO3→MO+CO2  (2)As an example of the above-mentioned scheme, FIG. 1 shows the variation in the free energy of the carbonation reaction as a function of temperature for calcium oxide. From the figure, we can see that the carbonation reaction is thermodynamically favored with a decrease in temperature (Gibbs free energy declines with a decrease in temperature). However, at lower temperatures, the carbonation reaction is kinetically slow. In fact, it takes geological time scales for the formation of CaCO3 by the reaction between CaO and atmospheric CO2 (at 280-360 ppm) at ambient temperatures. It should also be noted that the carbonation reaction would be favored as long as the free energy is negative. This creates an upper bound of 890° C. for carbonation to occur under a CO2 partial pressure of 1 atm. The equilibrium temperature for this reaction is a function of the partial pressure of CO2. A reaction based CO2 separation process offers many advantages. Under ideal conditions, MEA captures 60 g CO2/kg, silica gel adsorbs 13.2 g CO2/kg and activated carbon adsorbs 88 g CO2/kg. The sorption capacity of some metal oxides (such as the modified CaO, presented in this study) is about 700 g CO2/kg of CaO. This is about an order of magnitude higher than the capacity of adsorbents/solvents used in other CO2 separation processes and would significantly reduce the size of the reactors and the material handling associated with CO2 separation.
Numerous metal oxides exhibit the carbonation and calcination reaction. The calcination temperature of a few metal carbonates (CaCO3˜750° C., MgCO3˜385° C., ZnCO3˜340° C., PbCO3˜350° C., CuCO3˜225-290° C. and MnCO3˜440° C.) makes them viable candidates for this process. Apart from CaO, gas-solid carbonation of other metal oxides has not been widely studied. The carbonation of ZnO to ZnCO3 at 8-13° C. was low when exposed to CO2 and H2O for over 100 days (Sawada, Y.; Murakami, M.; Nishide, T. Thermal analysis of basic zinc carbonate. Part 1. Carbonation process of zinc oxide powders at 8 and 13° C. Thermochim. Acta. 1996, 273, 95-102.). MnCO3 undergoes a more complex thermal degradation phenomena. MnCO3 first decomposes to MnO2 at 300° C., which in turn changes to Mn2O3 at 440° C. At higher temperatures (˜900° C.), the final thermal decomposition product was identified as Mn3O4 (Shaheen, W. M.; Selim, M. M. Effect of thermal treatment on physicochemical properties of pure and mixed manganese carbonate and basic copper carbonate. Thermochim. Acta. 1998, 322(2), 117-128.). Different oxides of manganese provide the flexibility of exploiting the carbonation/calcination reaction over a wider temperature range. Aqueous phase MgO carbonation has been studied for its suitability for mineral-based CO2 sequestration (Fernandez, A. I.; Chimenos, J. M.; Segarra, M.; Fernandez, M. A.; Espiell, F. Kinetic study of carbonation of MgO slurries. Hydrometallurgy. 1999, 53, 155-167). The carbonation extent of Mg(OH)2 was about 10% between 387-400° C. and 6% formation between 475-500° C. (Butt, D. P.; Lackner, K. S.; Wendt, C. H.; Conzone, S. D.; Kung, H.; Lu, Y-C.; Bremser, J. K. Kinetics of Thermal Dehydroxylation and Carbonation of Magnesium Hydroxide. J. Am. Ceram. Soc. 1996, 79(7), 1892-1898). They attributed the low conversions to the formation of a non-porous carbonate product layer. This layer hinders the inward diffusion of CO2 and the outward diffusion of H2O (a product of the carbonation reaction) leading to low conversions. The carbonation of PbO was studied as a part of the chemical heat pump process (Kato, Y.; Saku, D.; Harada, N.; Yoshizawa, Y. Utilization of High Temperature Heat from Nuclear Reactor using Inorganic Chemical Heat Pump. Progress in Nuclear Energy. 1998, 32(3-4), 563-570. & Kato, Y.; Harada, N.; Yoshizawa, Y. Kinetic feasibility of a chemical heat pump for heat utilization from high temperature processes. Applied Thermal Engineering. 1999, 19, 239-254). They reported 30% conversion in an hour under 100% CO2 atmosphere at 300° C. Furthermore, they found the reactivity of PbO to drop with the number of carbonation-calcination cycles.
Carbonation of calcium oxide has been widely studied. Related applications of the CaO carbonation and calcination include the storage of energy (Barker, R. The Reversibility of the Reaction CaCO3=CaO+CO2. J. Appl. Chem. Biotechnol. 1973, 23, 733-742) and the zero emission coal alliance process, consisting of hydrogasification of coal fueled by the heat of the carbonation reaction (Tinkler, M. J.; Cheh, C. Towards a Coal-capable Solid Oxide Fuel Cell System. Proceedings of the 26th International Technical Conference on Coal Utilization and Fuel Systems; Clearwater, Fla., Mar. 5-8, 2001; pp 569-570). The gas-solid CaO—CO2 reaction proceeds through two rate-controlling regimes. The first regime involves a rapid, heterogeneous chemical reaction. In the second regime, the reaction slows down due to the formation of an impervious layer of CaCO3. This product layer prevents the exposure of unreacted CaO in the particle core to CO2 for further carbonation. The kinetics of the second regime is governed by the diffusion of ions through the CaCO3 product layer. The activation energy was estimated to be 21 kcal/mol below 688 K and 43 kcal/mol above it for the product layer diffusion, based on the counter migration of CO32− and O2− ions through the product layer (Bhatia, S. K.; and Perlmutter, D. D. Effect of the product layer on the kinetics of the CO2-Lime Reaction. AlChE J. 1983, 29(1), 79-86).
The extent of the carbonation reaction reported in many studies has also shown considerable variation. Stoichiometrically, 56 g of CaO should react with 44 g of CO2 to form 100 g of CaCO3. This translates to about 78.6-wt % capacity for CaO. However, the structural limitations prevent the attainment of theoretical conversion. The extent of carbonation was only 23-wt % in 30 minutes at 600° C. (Dedman, A. J.; Owen, A. J. Calcium Cyanamide Synthesis, Part 4.—The reaction CaO+CO2=CaCO3. Trans. Faraday Soc. 1962, 58, 2027-2035). A higher surface area CaO sorbent provided 55-wt % CO2 sorption (Bhatia, S. K.; and Perlmutter, D. D. Effect of the product layer on the kinetics of the CO2-Lime Reaction. AlChE J. 1983, 29(1), 79-86). 64-wt % CO2 sorption was achieved at 1050° C. temperature and 11.74 atm CO2 pressure in 32 hours (Mess, D.; Sarofim, A. F.; Longwell, J. P. Product Layer Diffusion during the Reaction of Calcium Oxide with Carbon Dioxide. Energy and Fuels. 1999, 13, 999-1005). However, the extent of carbonation at lower temperature/pressure conditions that are more characteristic of CO2 containing gaseous mixtures is absent in their work. The limitation in total conversion stems essentially from the nature of the initial pore size distribution of the CaO sorbent. Microporous sorbents (pore size<2 nm) are very susceptible to pore blockage and plugging due to the formation of higher molar volume product (molar volume of CaO: 17 cm3/mol; molar volume of CaCO3: 37 cm3/mol). CaO sorbents obtained from naturally occurring precursors are usually microporous in nature. At the end of the kinetically controlled regime, diffusion processes through the product layer control the reaction rate. Similar structural limitations have prevented calcium-based sorbents from attaining theoretical conversion for the sulfation reaction between CaO and sulfur dioxide (SO2) as well (Wei, S.-H.; Mahuli, S. K.; Agnihotri, R.; Fan, L.-S. High Surface Area Calcium Carbonate: Pore Structural Properties and Sulfation Characteristics. Ind. Eng. Chem. Res. 1997, 36(6), 2141-2148). They suggested that a mesoporous structure, which maximizes porosity in the 5-20 nm pore size range, would be less susceptible to pore pluggage. This structure would also be able to provide sufficient surface area to ensure rapid kinetics. Their modified precipitation technique resulted in a mesoporous CaCO3 structure that also had a high BET surface area determined by nitrogen (60 m2/g). A similar approach could also enhance the reactivity of CaO sorbents towards the carbonation reaction, which is the focus of this study.
Lastly, it is important that the CaO sorbents maintain their reactivity over many carbonation and calcination cycles. The conversion of CaO dropped from about 73% in the first carbonation cycle to 43% at the end of the 5th cycle at 866° C. (Barker, R. The Reversibility of the Reaction CaCO3=CaO+CO2. J. Appl. Chem. Biotechnol. 1973, 23, 733-742 & Barker, R. The Reactivity of Calcium Oxide Towards Carbon Dioxide and its use for Energy Storage. J. Appl. Chem. Biotechnol. 1974, 24, 221-227). Barker suggested that the CaCO3 layer is about 22 nm thick and his latter work showed repeated 93% conversion over 30 cycles at 629° C. on 10 nm CaO particles. In another study, cyclical studies conducted at a carbonation temperature of 880° C. and calcination at 860° C. led to a drop in conversion from 70% in the first carbonation to 38% in the 7th carbonation step (Kato, Y.; Harada, N.; Yoshizawa, Y. Kinetic feasibility of a chemical heat pump for heat utilization from high temperature processes. Applied Thermal Engineering. 1999, 19, 239-254). The process described here leads to >95% conversion due to the application of novel mesoporous CaO sorbents for CO2 capture and maintains their reactivity over repeated cycles of carbonation and calcination.